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The Components of Matter

Chapter 2. The Components of Matter. Categories of Matter. MATTER. MIXTURES. PURE SUBSTANCES. ELEMENTS. COMPOUNDS. HOMOGENOUS. HETEROGENEOUS. Composed only of atoms of the same element. Two or more elements chemically combined. (See definitions in slide after next.).

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The Components of Matter

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  1. Chapter 2 The Components of Matter

  2. Categories of Matter MATTER MIXTURES PURE SUBSTANCES ELEMENTS COMPOUNDS HOMOGENOUS HETEROGENEOUS Composed only of atoms of the same element. Two or more elements chemically combined. (See definitions in slide after next.)

  3. Definitions for Components of Matter Compound - a substance composed of two or more elements which are chemically combined. Figure 2.1 Mixture - a group of two or more elements and/or compounds which are physically intermingled.

  4. S Fe Physically mixed therefore can be separated by physical means; in this case by a magnet. Allowed to react chemically therefore cannot be separated by physical means. Figure 2.19 The distinction between mixtures and compounds. MIXTURE COMPOUND

  5. Mixtures Heterogeneous mixtures :have one or more visible boundaries between the components. Ex: water and oil Homogeneous mixtures: have no visible boundaries because the components are mixed as individual atoms, ions, and molecules. Ex: water and dissolved salt. Solutions :A homogeneous mixture is also called a solution. Solutions in water are called aqueous solutions, and are very important in chemistry. Although we normally think of solutions as liquids, they can exist in all three physical states.

  6. Tools of the Laboratory Basic Separation Techniques Filtration : Separates components of a mixture based upon differences in particle size. Normally separating a precipitate from a solution, or particles from an air stream. Crystallization : Separation is based upon differences in solubility of components in a mixture. Distillation : separation is based upon differences in volatility. Extraction : Separation is based upon differences in solubilityin different solvents (major material). Chromatography : Separation is based upon differences in solubilityin a solvent versus a stationary phase.

  7. Basic Separation Techniques From full-size Silberberg text Figure B2.3 Filtration Figure B2.4 Crystallization

  8. PROPERTIES OF MATTER Physical Properties: color, texture, odor, density, BP, MP, etc. Chemical Properties: how a substance reacts with other substances

  9. Chemical Physical Properties*: Physical properties: MP, BP, color, density, solubility in water. (*Except Na reacts with water.)

  10. Be able to: • Identify the type of matter in four ways: element, compound, homogenous mixture, or heterogeneous mixture • Know the three states of matter • Separate physical properties from chemical properties

  11. total mass total mass calcium oxide + carbon dioxide calcium carbonate CaO + CO2 CaCO3 Law of Mass Conservation: The total mass of substances does not change during a chemical reaction. reactant 1 + reactant 2 product = 56.08g + 44.00g 100.08g

  12. 8.0 g calcium 2.4 g carbon 9.6 g oxygen 0.40 calcium 0.12 carbon 0.48 oxygen 40% calcium 12% carbon 48% oxygen 20.0 g 1.00 part by mass 100% by mass Figure 2.3 (4th ed.) Law of Definite (or Constant) Composition: No matter the source, a particular compound is composed of the same elements in the same parts (fractions) by mass. Calcium carbonate Analysis by Mass (grams/20.0g) Mass Fraction (parts/1.00 part) Percent by Mass (parts/100 parts)

  13. g O g O 72.7 57.1 g C g C = 27.3 42.9 = 1.33 2.66 g O/g C in II 2 = 1.33 g O/g C in I 1 = 2.66 Law of Multiple Proportions: If elements A and B react to form two compounds, the different masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers. Example: There are two “Carbon Oxides,” A & B Carbon Oxide I : 57.1% oxygen and 42.9% carbon Carbon Oxide II : 72.7% oxygen and 27.3% carbon Assume that you have 100 g of each compound. That means that Oxide I has 57.1 g of oxygen and 42. 9 g of carbon; Oxide II has 72.7 g of oxygen and 27.3 g of carbon. =

  14. Dalton’s Atomic Theory (MEMORIZE!) 1. All matter consists of atoms. 2. Atoms of one element cannot be converted into atoms of another element. 3. Atoms of an element are identical in mass and other properties and are different from atoms of any other element. 4. Compounds result from the chemical combination of a specific ratio of atoms of different elements.

  15. Fig 2.4 Experiments to Determine the Properties of Cathode Rays CONCLUSION OBSERVATION 1. Ray bends in magnetic field. 2. Ray bends towards positive plate in electric field. consists of charged particles consists of negative particles 3. Ray is identical for any cathode. This led to the discovery of Electrons. particles found in all matter

  16. Millikan’s oil-drop experiment for measuring an electron’s charge. Figure 2.5 (1909)

  17. determined by J.J. Thomson and others mass charge mass of electron = X charge Millikan used his findings to also calculate the mass of an electron. = (-5.686x10-12 kg/C) X (-1.602x10-19 C) = 9.109x10-31 kg = 9.109x10-28 g

  18. Early Atomic Theory At the beginning of the 20th Century: scientists knew an atom contained both positive and negative particles. One model was “plum pudding” – like raisins in rice pudding. Rutherford’s experiment showed that an alpha particle (a Helium nucleus with 2 protons and 2 neutrons) aimed at gold foil was deflected almost straight back. NOT WHAT HE EXPECTED: HAD TO REVISE THEIR MODEL!!! Led to modern nuclear atom model.

  19. Figure 2.6

  20. Figure 2.7 General features of the nuclear atom (Know these!) Atom is electrically neutral, spherical & composed of a positively charged central nucleus surrounded by one or more negatively charge electrons. Nucleus consists of protons and neutrons.

  21. Location in the Atom Name(Symbol) Relative Absolute(C)* Relative(amu)† Absolute(g) Proton (p+) 1+ +1.60218x10-19 1.00727 1.67262x10-24 Nucleus Neutron (n0) 0 0 1.00866 1.67493x10-24 Nucleus Outside Nucleus Electron (e-) 1- -1.60218x10-19 0.00054858 9.10939x10-28 Properties of the Three Key Subatomic Particles Table 2.2 Charge Mass You memorize what’s boxed above. * The coulomb (C) is the SI unit of charge. † The atomic mass unit (amu) equals 1.66054x10-24 g.

  22. Atomic Symbols, Isotopes, Numbers A X The Symbol of the Atom or Isotope Z X = Atomic symbol of the element A = mass number; A = Z + N Z = atomic number (the number of protons in the nucleus) N = number of neutrons in the nucleus Isotope = atoms of an element with the same number of protons, but a different number of neutrons From Figure 2.8

  23. PROBLEM: Silicon has three naturally occurring isotopes: 28Si, 29Si, and 30Si. Determine the number of protons, neutrons, and electrons in each silicon isotope. Sample Problem 2.4 Determining the Number of Subatomic Particles in the Isotopes of an Element The atomic number of silicon is 14. Therefore SOLUTION: 28Si has 14p+, 14e- and 14n0 (28-14) 29Si has 14p+, 14e- and 15n0 (29-14) 30Si has 14p+, 14e- and 16n0 (30-14) Now you try these uranium isotopes: 234, 235, 238, and 239.

  24. ATOMS & ISOTOPES REMEMBER: In an atom of the same element # protons never changes (except for nuclear decay) # neutrons can be different  ISOTOPES # electrons can change  IONS

  25. The Mass Spectrometer and Its Data Tools of the Laboratory Figure B2.2 2.9

  26. AVERAGE ATOMIC MASS Fig 2.9 is a simple schematic, but the mass spec determines the actual mass of each isotope to many sig figs and its percent abundance, usually to six or more sig figs. Chart B = the mass charge ratio, which determines the mass to many sig figs (look at fluorine on per table) Chart C = a count of each isotope which is reported as percent abundance, again to many sig figs By using a mass spectrometer, an atom’s isotopes can be counted, establishing relative abundance of each isotope and its exact mass. From this data we can calculate a weighted average atomic mass: Weighted average atomic mass = fract1*mass1 + fract2*mass2 + ...

  27. Example of calculating average atomic mass We put boron in a mass spectrometer, and find there are just two isotopes, with these results: 19.91% Boron-10 with mass 10.0129 amu and 80.09% Boron-11 mass with11.0093 amu Average atomic mass = (0.1991 * 10.0129) + (0.8009 * 11.0093) = 1.994 + 8.817 = 10.811 (sig fig rules) Do problem 33 for practice on this method. (Magnesium has three naturally occurring isotopes. 24Mg has a mass of 23.9850 amu and 78.99% abundance, 25Mg has a mass of 24.9868 amu and 10.00% abundance, and 26Mg has a mass of 25.9826 amu and 11.01% abundance.)

  28. Difference between atomic mass and mass number You have just practiced determining the weighted average atomic mass. How is that different from mass number? ATOMIC MASS IS NOT THE SAME AS MASS NUMBER! ATOMIC MASS IS NOT THE SAME AS MASS NUMBER! ATOMIC MASS IS NOT THE SAME AS MASS NUMBER! Think you can remember this?

  29. The Modern Reassessment of the Atomic Theory 1. All matter is composed of atoms. The atom is the smallest body thatretains the unique identityof the element. 2. Atoms of one element cannot be converted into atoms of another element in a chemical reaction. Elements can only be converted into other elements in nuclear reactions. 3. All atoms of an element have the same number of protons and electrons, which determines the chemical behavior of the element. Isotopes of an element differ in the number of neutrons, and thus in mass number. A sample of the element is treated as though its atoms have anaverage mass. 4. Compounds are formed by the chemical combination of two or more elements in specific ratios.

  30. Figure 2.10 The modern periodic table.

  31. Tell me what you know about groups, columns, trends, etc. Write down group names, locations for metal, nonmetals and metalloids. (Alkali metals, alkaline earth metals, halogens, noble gases, metalloids along stairs, metals to lower left, nonmetals to upper right) What trends do you know? (Fr is largest, He is smallest) Where are the transition metals and the inner transition metals? (Be able to locate them.)

  32. You need to memorize the first 36 elements’ symbols and names, plus ten more: Ag, Sn, I, Ba, Pt, Au, Hg, Pb, Rn, and U.

  33. Figure 2.11 from 4th ed. Metals, metalloids, and nonmetals. Cadmium Copper Lead Chromium Bismuth Arsenic Chlorine Silicon Antimony Bromine Sulfur Iodine Carbon (graphite) Boron Tellurium Stop here and go to chapter 7!

  34. Figure 2.14 from 4th ed. The relationship between ions formed and the nearest noble gas.

  35. Formation of a covalent bond between two H atoms. Figure 2.13 Covalent bonds form when elements share electrons, which usually occurs between nonmetals.

  36. Section 2.8 – after chp 11 • Use the lecture notes, the textbook AND the lab manual’s Dry Lab on Nomenclature. • All the tables in the dry lab will be very useful!

  37. You must memorize these! The seven common diatomic elements: H2, O2, N2, F2, Cl2, Br2, I2 (Memory aid: HON+Halogens)

  38. Types of Chemical Formulas A chemical formula is comprised of element symbols and numerical subscripts that show the type and number of each atom present in the smallest unit of the substance. An empirical formula indicates the relative number of atoms of each element in the compound. It is the simplest type of formula. The empirical formula for hydrogen peroxide is HO. A molecular formula shows the actual number of atoms of each element in a molecule of the compound. The molecular formula for hydrogen peroxide is H2O2. A structural formula shows the number of atoms and the bonds between them, that is, the relative placement and connections of atoms in the molecule. The structural formula for hydrogen peroxide is H-O-O-H.

  39. Some common monatomic ions of the elements. Figure 2.16 Can you see any patterns? See large Periodic Table handout for variety of charges of metals. Be2+

  40. Naming metal ions with more than one oxidation number (charge) Remember that metals like iron or tin with more than one possible charge have to show which one by using Roman numerals for the charge in parentheses. Sn2+ is tin (II) ion Sn4+ is tin (IV) ion

  41. Naming binary ionic compounds: see rules in your packet The name of the cation is written first, followed by that of the anion. The name of the cation is the same as the name of the metal. Many metal names end in -ium. The name of the anion takes the root of the nonmetal name and adds the suffix -ide. Calcium ion and bromide ion form calcium bromide.

  42. Naming Binary Ionic Compounds EXAMPLE: Form an ionic compound of Al & Br and then Al & O. Al forms Al3+ and Br forms Br-, so the ratio will be 1:3 or AlBr3 aluminum bromide. Again, Al forms Al3+ and O forms O2-, so the ratio will be 2:3 or Al2O3 aluminum oxide. Special trick called the criss-cross rule: Xa+ & Yb- XbYa but if b=a, reduce formula to 1:1 ratio (unless mercury (I) ion is involved)

  43. PROBLEM: Write formulas and name the ionic compound formed from the following pairs of elements: Naming Binary Ionic Compounds (sample problems 2.7 & 2.8 (a) magnesium and nitrogen (b) iodine and cadmium (c) strontium and fluorine (d) sulfur and cesium SOLUTION: (a) Mg2+ & N3-; three Mg2+(6+) & two N3-(6-); Mg3N2 magnesium nitride (b) Cd2+ & I-; one Cd2+(2+) & two I-(2-); CdI2 cadmium iodide (c) Sr2+ & F-; one Sr2+(2+) & two F-(2-); SrF2 strontium fluoride (d) Cs+ & S2-; two Cs+(2+) and one S2-(2-); Cs2S cesium sulfide Now do the follow-up problems.

  44. PROBLEM: Give the systematic names for the formulas or the formulas for the names of the following compounds: Sample Problem 2.9 Determining Names and Formulas of Ionic Compounds of Elements That Form More Than One Ion (a) tin(II) fluoride (b) CrI3 (d) CoS (c) Iron (III) oxide SOLUTION: (a) Tin (II) is Sn2+; fluoride is F-; so the formula is SnF2. (b) The anion I is iodide(I-); 3I- means that Cr(chromium) is +3. CrI3 is chromium(III) iodide (c) Iron (III) is the name for Fe3+; oxide is O2-, therefore the formula is Fe2O3. (d) Co is cobalt; the anion S is sulfide(2-) so cobalt must be +2; the compound is cobalt (II) sulfide.

  45. Hydrates of ionic compounds Hydrates are ionic crystals of salts with water molecules incorporated in their crystal structures. Write "formula unit name - dash - Greek prefix (representing # of water molecules) hydrate" BaCl2.2H2O is barium chloride dihydrate You try: Name CuSO4.5H2O: Give the formula for sodium sulfate decahydrate:

  46. per root ate ClO4- perchlorate root ate ClO3- chlorate No. of O atoms root ite ClO2- chlorite hypo root ite ClO- hypochlorite Prefix Prefix Prefix Number Number Number 1 mono 4 tetra 8 octa 2 di 5 penta 9 nona 3 tri 6 hexa 10 deca 7 hepta Naming oxoanions Figure 2.17 Root Suffixes Examples Prefixes Table 2.6 Numerical Prefixes for Hydrates and Binary Covalent Compounds

  47. PROBLEM: Give the systematic names or the formula or the formulas for the names of the following compounds: (c) Ba(OH)2 8H2O Sample Problem 2.10 Determining Names and Formulas of Ionic Compounds Containing Polyatomic Ions (a) Fe(ClO4)2 (b) sodium sulfite SOLUTION: (a) ClO4- is perchlorate; iron must have a 2+ charge. This is iron(II) perchlorate. (b) The anion sulfite is SO32- therefore you need 2 sodiums per sulfite. The formula is Na2SO3. (c) Hydroxide is OH- and barium is a 2+ ion. When water is included in the formula, we use the term “hydrate” and a prefix which indicates the number of waters. So it is barium hydroxide octahydrate.

  48. Naming Inorganic Acids 1) Binary acid solutions form when certain gaseous compounds dissolve in water. For example, when gaseous hydrogen chloride (HCl) dissolves in water, it forms a solution called hydrochloric acid. Prefix hydro- + anion nonmetal root + suffix -ic + the word acid - hydrochloric acid Practice naming H2S(aq): 2) Oxoacid names are similar to those of the oxoanions, except for two suffix changes: Anion “-ate” suffix becomes an “-ic” suffix in the acid. Anion “-ite” suffix becomes an “-ous” suffix in the acid. The oxoanion prefixes “hypo-” and “per-” are retained. Thus, BrO4- is perbromate, and HBrO4 is perbromic acid; IO2- is iodite, and HIO2 is iodous acid. Practice naming HClO3:

  49. PROBLEM: Name the following anions and give the names and formulas of the acids derived from them: Sample Problem 2.12 Determining Names and Formulas of Anions and Acids (a) Br - (b)IO3 - (c) CN - (d) SO4 2- (e) NO2 - SOLUTION: (a) The anion is bromide; the acid is hydrobromic acid, HBr. (b) The anion is iodate; the acid is iodic acid, HIO3. (c) The anion is cyanide; the acid is hydrocyanic acid, HCN. (d) The anion is sulfate; the acid is sulfuric acid, H2SO4. (e) The anion is nitrite; the acid is nitrous acid, HNO2.

  50. NAMING BINARY COVALENT COMPOUNDS Very simple – you will know a compound is covalent if it's two nonmetals. Indicate how many of each atom in the compounds by using Greek prefix (mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca) If first element is single, leave off mono. If first element is hydrogen, leave off any prefix. ‘Prefix'element name - 'prefix'element root - suffix 'ide' Practice: CO, CO2, NO2, N2O4, P2O5, HF, H2S

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